Figure 11-2. The Activity Series for Metals

In this list, metals are arranged according to their ability to lose electrons (be oxidized). The more easily a metal is oxidized, the more reactive it tends to be. Indeed, the metals at the top of the list are highly reactive, while the metals at the bottom are known and used for their lack of reactivity. (Unit 11, Section 2)

Sidebar Image. A Solution of Potassium Permanganate

Potassium permanganate can oxidize many chemical compounds, which in turn reduces the manganese from its +7 oxidation state to a lower oxidation state. When this happens, the deep purple color disappears—a very noticeable indicator that a reaction has taken place. (Unit 11, Section 3)

Figure 11-4. Table of Standard Reduction Potentials

The redox reactivity of different chemical species is often tabulated according to standard reduction potentials for half-reactions. Any combination of half-reactions can, in principle, produce a spontaneous redox reaction; the direction of the reaction is determined by which half-reaction has the greater standard reduction potential. (Unit 11, Section 4)

Figure 11-7. Batteries

A selection of commercially available batteries for different personal-use applications is shown above. Many of these may look familiar as the power source for household items. Most of the ones pictured here are traditional alkaline batteries. (Unit 11, Section 5)

Figure 11-6. The Classic Galvanic Cell

The Cu/Zn reaction is set up in separate beakers with a salt bridge and wire connecting the two half-cells. The direction of electron flow is given above the wire, and the voltmeter displays a positive value for the cell emf, indicating a spontaneous reaction. (Unit 11, Section 5)

Figure 11-8. 9-Volt Battery

Batteries come in all shapes, sizes, and chemistries. In addition to adjusting the chemistry to produce a battery with a particular voltage, individual cells can be connected in a series to generate a larger voltage, as in the 9V alkaline battery. (Unit 11, Section 6)

Figure 11-9. Statue of Liberty

Pure copper oxidizes slowly, but copper objects left outside develop a greenish layer, called a "patina," over time. The exact composition of the patina depends on the environment to which the copper is exposed, but the resulting layer prevents the copper underneath from being oxidized further. (Unit 11, Section 7)

Figure 11-10. Two Anodized Aluminum Carabiners

First, the surface of the aluminum was oxidized to create a dull coating of aluminum oxide, which then allows the surface to be dyed an array of colors. Thus, it protects the interior aluminum metal from further oxidation and provides an aesthetic surface. (Unit 11, Section 7)

Sidebar Image. Galvanized Nail

Electroplating can be used to put a zinc coating, called galvanization, on the surface of an iron nail. This coating both protects the iron inside the nail and serves as a sacrificial anode oxidizing before the iron can, increasing the lifespan of the nail. (Unit 11, Section 7)

Figure 11-11. The Structure of Cisplatin

The central platinum metal atom has four coordinate covalent bonds to two chloride ions and two ammonia molecules, making a coordination complex that has a square planar geometry. Because this molecule is flat, the two chlorines could be next to each other or they could be across from each other on the metal. However, if they were across from each other, that would make a different isomer called "trans-platin." (Unit 11, Section 8)

Figure 11-12. Copper(II) Sulfate Pentahydrate

The copper ions in copper(II) sulfate pentahydrate are actually in an octahedral arrangement. There are six ligands that share their electrons with the copper metal through their oxygen atoms. The complex ion that forms is what gives copper(II) its blue color. (Unit 11, Section 8)

Figure 11-13. Ligand Field Splitting

The orbitals in a d subshell are usually all the same energy, but in transition metal compounds, the energy levels split. Electrons can be promoted from the lower energy levels to higher energy levels when the compound absorbs light of an appropriate energy, delta octahedral (Δo) in this case. (Unit 11, Section 9)

Figure 11-14. Coordination Complexes of Nickel

From left to right: [Ni(NH3)6]2+, [Ni(en)3]2+, [NiCl4]2-, [Ni(H2O)6]2+. Note how changing the number of ligands and the identity of the ligands on nickel(II) ions can create an array of different colors, as the energies of the possible electronic transitions of the d electrons are affected by the changes in the ligands. (Unit 11, Section 9)

Figure 11-15. A Color Wheel

The colors in the visible spectrum are absorbed by a colored object, and the rest of the visible light is reflected. This light is absorbed because of the possible electronic transitions in the compounds. If a compound absorbs green visible light, the remainder of the visible spectrum heads toward the eye, and the brain interprets this as the color opposite it on the color wheel. So, a compound that appears red to the eye absorbs green light. (Unit 11, Section 9)

Figure 11-16. Hemoglobin

The heme group in hemoglobin bonds to oxygen in the bloodstream. The oxygenated and deoxygenated forms of hemoglobin have different colors, and so the oxygen content in blood can be measured spectroscopically using a pulse oximeter. (Unit 11, Section 10)